Metallic Materials

Corrosion and Electrochemical Corrosion Protection

2.2 Electrochemical Corrosion

2.2.1 Metallic Materials

Metallic materials consist of one or more metallic phases, depending on their composition, and very small amounts of nonmetallic inclusions. In the metallic state, atoms donate some of their outer electrons to the electron gas that permeates the entire volume of the metal and is responsible for good electrical conductivity (10⁵ S cm"¹). Pure elements do not react electrochemically as a single component.

A mesomeric state can be approximately assumed

in which iron ions and electrons appear in the metal. Both components can react with electrolytes in which the dissolution of metal by the passage of Fe²⁺ results in a positive current 7A and metal loss of Am, while the passage of electrons corre- sponds to a negative current 7_C without removal of metal. In the first case an anodic reaction occurs, and in the second a cathodic reaction. The relations are reversed in

a transfer from the solution to the metal: in the anodic reaction, electrons are trans- ferred to the metal, and in the cathodic reaction iron is deposited. Cathodic metal deposition is used in electroplating and is the reverse process of electrolytic corro- sion. Faraday's Law applies to both:

where Am is the mass of dissolved metal, M is the atomic weight, Q is the transferred electric charge, z is the valence of the metal ions, and & is Faraday's constant. The following equations can be derived with the aid of the specific gravity ps:

where JA is current density of the anodic partial reaction for the passage of metal ions; S is the surface area of the electrode; t is time; and As is the thickness of material removed. Equation (2-4) can be divided as follows:

where w is the rate of decrease in thickness and v is the rate of weight loss per unit surface area. The conversion factors fa and/^ for some important metals are given in Table 2-1. A further relationship is given by:

These corrosion parameters have to be modified for time- and place-related reaction velocities [6]. Different local removal rates are in general due to differ- ences in composition or nonuniform surface films, where both thermodynamic and

Table 2-1 Conversion factors and standard potentials for electrochemical metal- metal ion reactions

Reaction from Eq. (2-21)

Ag = Ag⁺ + e- 2 Hg = Hg²⁺ + 2 e- Cu = Cu²⁺ + 2 e- Pb = Pb²⁺ + 2 e- Mo = Mo³⁺ + 3 e- Ni = Ni²⁺ + 2 e- Tl = Tl⁺ + e- Cd = Cd²⁺ + e- Fe = Fe²⁺ + 2 e~

Cr = Cr³⁺ + 3 e~

Zn = Zn²⁺ + 2 e- Cr = Cr²⁺ + 2 e- Mn = Mn²⁺ + 2 e~

Al=Al^{3 +} + 3e- Mg = Mg²⁺ + 2 e-

/.

Eq. (2-5) mm a"¹ mA cm~²

33.6 - 11.6 29.8 10.2 10.8 56.4 21.2 11.6 8.2 15.0 12.3 12.5 10.9 22.8

fb

Eq. (2-6) gm-'h-¹ mA cm"²

40.2 - 11.9 38.6 11.9 11.0 76.2 21.0 10.4 6.5 12.2 9.7 10.2 3.35 4.54

fc

Eq. (2-7) mm a"¹ grn^h'¹

0.83 - 0.98 0.77 0.86 0.98 0.74 1.01 1.12 1.27 1.23 1.27 1.22 3.24 5.03

I/a (25°C) Eq. (2-29)

V +0.80 +0.80 +0.34 -0.13 -0.20 -0.24 -0.34 -0.40 -0.44 -0.74 -0.76 -0.91 -1.18 -1.66 -2.38 kinetic effects have an influence. The tendency of metallic materials to local corro- sion can be characterized as follows:

(a) uniform weight loss occurs mostly on active, almost single-phase metals;

(b) shallow pitting and pitting in general are only possible in the presence of surface films, particularly on passive metals;

(c) selective corrosion is only possible on multiphase alloys.

There are no films or protective surface films on active metals, e.g., mild steel in acid or saline solutions. Passive metals are protected by dense, less readily soluble surface films (see Section 2.3.1.2). These include, for example, high-alloy Cr steels and NiCr alloys as well as Al and Ti in neutral solutions. Selective corrosion of alloys is largely a result of local concentration differences of alloying elements which are important for corrosion resistance e.g., Cr [4].

The passage of electrons from the metal to the electrolyte is not directly related to metal removal, but has an indirect connection due to the electron neutrality law:

Electrons cannot be dissolved in aqueous solutions but react with oxidants in the following way:

Ox and Red are general symbols for oxidation and reduction media respectively, and n and (n-z) indicate their numerical charge (see Section 2.2.2). Where there is no electrochemical redox reaction [Eq. (2-9)], the corrosion rate according to Eq. (2-4) is zero because of Eq. (2-8). This is roughly the case with passive metals whose surface films are electrical insulators (e.g., Al and Ti). Equation (2-8) does not take into account the possibility of electrons being diverted through a conduc- tor. In this case the equilibrium

is valid instead of Eq. (2-8).

The current 7 is called the total current. In free corrosion, i.e., without the contribution of external currents (see Fig. 2-1), it is always zero, as given by Eq. (2- 8). 7_A and 7_C are known as the anodic and cathodic partial currents. According to Eq. (2-10), generally in electrolytic corrosion anodic total currents and/or cathodic redox reactions are responsible.

All metallic materials can suffer electrolytic corrosion. Fractures caused by cathodic hydrogen only occur when the activity of the absorbed hydrogen and the level of the tensile stress, which can be external or internal, reach a critical value.

In general, critical hydrogen absorption is achieved only in the presence of pro- moters. However, under very severe conditions such as at very low pH or very negative potential, critical hydrogen absorption can occur. Steels with a hardness greater than HV 350 are particularly susceptible.

Materials consisting of elements of subgroups 4 and 5 of the periodic table are prone to the formation of internal hydrides, leading to severe embrittlement and fracture. Titanium and zirconium are important examples. Materials consisting of elements in the main groups 4 to 6 of the periodic table suffer weight loss by corrosion due to the formation of volatile hydrides [7]. A typical example is lead.

Types of corrosion arising from cathodic hydrogen can limit the application of cathodic protection and are dealt with in Refs. 8 and 9.

2.2.2 Aqueous Electrolytes

Anions and cations exist in water. They migrate in an electric field and thus carry a current. Ohm's Law is applicable:

where % is the specific conductivity and E is the electric field strength. In dilute electrolytes, the conductivity is the sum of the ion mobilities /(:

The index i represents the type of ion and c is its concentration. In water, the ions have velocity wF;, giving the relation:

The quotient u{ is called the electrochemical mobility and is tabulated along with ion mobilities. It is important to pay attention to the units because of possible confusion. Values of /( are given in Table 2-2. Raising the temperature usually in- creases ion mobility, while increasing the concentration reduces the conductivity due to interactions:

Electrical conductivity is of interest in corrosion processes in cell formation (see Section 2.2.4.2), in stray currents, and in electrochemical protection methods.

Conductivity is increased by dissolved salts even though they do not take part in the corrosion process. Similarly, the corrosion rate of carbon steels in brine, which is influenced by oxygen content according to Eq. (2-9), is not affected by the salt concentration [4]. Nevertheless, dissolved salts have a strong indirect influence on many local corrosion processes. For instance, chloride ions that accumulate at lo- cal anodes can stimulate dissolution of iron and prevent the formation of a film.

Alkali ions are usually regarded as completely harmless, but as counterions to OH~

ions in cathodic regions, they result in very high pH values and aid formation of films (see Section 2.2.4.2 and Chapter 4).

The pH value and thus the OH~ ion concentration is important in the formation of surface films, since OH~ ions generally form difficultly soluble compounds with metal ions (see Section 2.2.3.1). pH is an important parameter of the medium. One

Table 2-2 Ion mobilities /. in S cm² mol~' for calculating specific conductivity with Eq. (2-12); between 10 and 25 °C, conductivity increases between 2 and 3% per °C

£ r - 96,487 —J x- S cm² mol"¹ V^-1 s'¹ cm²

Cation +z / at Anion -z / at_{; i}

25°C 100°C 25°C 100°C

H3O⁺ 1 350 637 OH 1 200 446

Na⁺ 1 50 150 Cl- 1 76 207

K⁺ 1 73 200 NO' 1 71 189

Mg²⁺ 2 53 170 HCO- 1 44

Ca²⁺ 2 59 187 CO²" 2 72 -

Fe²⁺ 2 53 S0²~ 2 79 256

Cu²⁺ 2 56 -

has to remember, however, that considerable changes in the pH value can occur as a result of subsequent reactions on the metal surface. Generally the equilibrium is:

(Deviations from the ideal behavior that could be accounted for by the use of activ- ity coefficients are neglected here.)

The concentration of oxidizing agents is essential for the course of reactions involving Eq. (2-9). These can be divided into two groups according to the type of oxidizing agent:

(a) oxygen corrosion (in all media)

(b) acid corrosion (particularly in strong acids)

The evolution of hydrogen from the acid molecule can also occur in slightly dissociated weak acids such as H2CO3 and H2S. In the case of only slightly dissoci- ated weak acids, such as H2CO3 and H2S, production of hydrogen can also occur from the acid molecules. In this case, the acid concentration rather than the pH value is a measure of the aggressiveness of the corrosion. In the same way, hydro- gen can be evolved from H2O:

This reaction occurs with overall cathodic currents, i.e., with cathodic polariza- tion. It can be practically ignored in the case of free corrosion of steel in a neutral solution. Other oxidizing media are of interest only in special cases.

In some cases the amount of gas evolved or consumed according to Eqs. (2-17) to (2-19) can be asked for. Then Eq. (2-6) is applicable. To obtain the volume, the following applies:

where jv is the rate of evolution or consumption of gas; Vis the volume of gas and V is the mol volume under standard conditions. For V = 22.4 L it follows that:

Some data are given in Table 2-3.

2.2.3 Electrochemical Phase Boundary Reactions

In electrolytic corrosion, an anodic partial reaction takes place according to Eq. (2-3)

and a cathodic redox reaction according to Eq. (2-9) (see Fig. 2-1). The reaction rates can be expressed in general by using Eq. (2-6) with equivalent currents 7A and 7C. They are a function of the partners to the reaction and the potential U. For every partial reaction there is an equilibrium potential if in which the overall reaction is zero. The following section deals with the thermodynamic and kinetic fundamen- tals of these reactions.

Table 2-3 Conversion factors and standard potentials for electrochemical redox reactions

1/&(25°C) f_b /„

Eq. (2-30) Eq. (2-6) Eq. (2-20) Reaction according to g m"² h"' L m~ h~

Eq. (2-9) V mAcnT² mA cm"²

2H⁺ + 2e~ = H2 0.00 0.37 4.18

O2 + 2H2O + 4e- = 4OH- +0.40 2.98 2.09

Cl2 + 2e- = 2Cl- +1.36 13.24 4.18

Cr²⁺ + e- = Cr⁺ -0.41

Cu²⁺ + e- = Cu⁺ +0.16

Fe²⁺ + e- = Fe²⁺ +0.77

2.2.3.1 Basic Thermodynamics

The driving force for the transport of all particles is a change in the electro- chemical potential/I( which is related to the partial molar free enthalpy/t, and the electric potential 0 as follows:

For a homogeneous conductor and in the migration direction [see Eq. (3-1)]:

The factor B = DIRT is the mobility and contains the diffusion coefficient D, the gas constant R, and the absolute temperature T. The equation includes a diffusion and a migration term. Correspondingly Eq. (2-23) gives the first diffusion law for Zi = 0 and Ohm's Law for grad fa = 0. For transfer across a phase boundary:

Finally if \v_i = 0, the equilibrium is given by:

Application of Eq. (2-25) to the reaction under consideration, i.e., Eq. (2-21), and to the potential-determining reaction of the reference electrode, Eq. (2-18), leads with Eq. (2-1) to the Nernst potential equation:

In this example, AG is the free reaction enthalpy of the chemical reaction

which corresponds to adding the electrochemical reactions in Eq. (2-18) and (2-21).

The negative sign of If accounts for the fact that all A0 contain potential differences in the reaction direction of Eq. (2-27) in the cell H2/electrolyte/metal and AG is appropriately defined [10]. From the concentration dependence of/I, it follows that

and for the standard state of the hydrogen electrode with a variable metal ion concentration c(Me^z+), the equilibrium potential against the standard hydrogen elec- trode is

and U£ is the standard potential which can be calculated from the free enthalpy of formation AG° Table 2-1 shows the more important values. The factor RT/ZP is 26 mV at 25 °C. In the same way, a potential equilibrium can be derived for a simple redox reaction [Eq. (2-9)]:

Many redox reactions are more complicated than that given by Eq. (2-9). For a general redox reaction, with components Xt and their coefficients n{ written as

i

the relation can be derived [4,10]:

A comprehensive list of standard potentials is found in Ref. 7. Table 2-3 gives a few values for redox reactions. Since most metal ions react with OH" ions to form solid corrosion products giving protective surface films, it is appropriate to represent the corrosion behavior of metals in aqueous solutions in terms of pH and UH. Figure 2-2 shows a Pourbaix diagram for the system Fe/H2O. The boundary lines correspond to the equilibria:

line (1): Fe/Fe²⁺ corresponding to Eq. (2-21),

line (2): Fe/Fe(OH)₂ from Fe + 2 H₂O = Fe(OH)₂ + 2 H⁺ + 2 e- line (3): Fe²⁺/FeOOH from Fe²⁺ + 2 H2O = FeOOH + 3 H⁺ + e~

line (4): Fe³⁺/FeOOH from Fe³⁺ + 2 H₂O - FeOOH + 3 H⁺, line (a): OH7O₂ corresponding to Eq. (2-17),

line (b): H₂/H⁺ corresponding to Eq. (2-18).

Electrolytic corrosion occurs in regions I and IV with the formation of soluble iron ions. Solid corrosion products which can have a protective effect are formed in region II. This is the region of surface film passivity. Certain corrosive sub-

Fig. 2-2 Simplified potential- pH diagram for an iron/aqueous electrolyte system at 25°C;

c(Fe²⁺) + c(Fe³⁺) = 1Q-⁶ mol L"¹ (explanation in the text).

stances in the medium (e.g., chloride ions) and mechanical effects can destroy surface films locally, leading to intensive local corrosion such as pitting and stress corrosion. On the other hand, in certain passivating acids, such as HNO3, H2SO4, and H3PO4, there is an area in region I (shown hatched), Fig. 2-2, where the mate- rial is covered with an extremely thin nonequilibrium film. This chemical passivity is not technically different from surface film passivity. For both cases, corrosion rates are extremely small but not zero, as in region III where the metal is thermo- dynamically stable. In addition, there is the latent danger of local corrosion [4].

Such pH-£/H diagrams are available for all metals [7]. They give an overview of corrosion behavior and the electrochemical protection that is possible when the potential is changed by means of impressed currents. Cathodic currents are required to reduce the potential in region III. Anodic currents are needed to raise the potential in region II or the hatched area in region I. This is the basis of cathodic and anodic protection. The region of H2O stability between the straight lines (a) and (b) must be considered before a preliminary judgment can be made. Outside these lines, the possibility of changing the potential is limited by the electrolytic dissociation of water. From Fig. 2-2 it follows that cathodic protection in acid solutions is not prac- tically possible but on the other hand anodic protection probably is possible.

Regions for soluble hydroxy complexes of the type Fe(OH)⁺ are neglected in Fig. 2-2 and in corresponding diagrams in Ref. 7; compare the amended diagrams in the literature [11,12]. The regions in the pH-potential diagrams can be transposed to some extent by complex formation. This must be taken into account when dealing with unknown chemical solutions. In addition, when using the pH-potential diagrams in Ref. 7, care must be taken that, although regions of weight loss due to hydride formation are represented, regions where internal hydride formation occurs in met- als of the 4 and 5 subgroups are not given.

2.2.3.2 Electrochemical Kinetics

The potential dependence of the velocity of an electrochemical phase bound- ary reaction is represented by a current-potential curve /(£/)• It is convenient to relate such curves to the geometric electrode surface area S, i.e., to present them as current-density-potential curves J(U). The determination of such curves is repre- sented schematically in Fig. 2-3. A current is conducted to the counterelectrode E3

in the electrolyte by means of an external circuit (voltage source U0, ammeter, resistances R' and R"} and via the electrode E} to be measured, back to the external circuit. In the diagram, the current indicated (®) is positive. The potential of E} is measured with a high-resistance voltmeter as the voltage difference of electrodes Ej and E2. To accomplish this, the reference electrode, E2, must be equipped with a Haber-Luggin capillary whose probe end must be brought as close as possible to

Fig. 2-3 Potential scheme (a) and circuit (b) for measuring a current-potential curve in cathodic polarization (explanation in the text).

the surface Ej without shielding the current flow [13]. The potential diagram shows that the current in the electrolyte produces a voltage drop rfo that basically always results in an error in the potential measurement:

By comparison with Eq. (2-1) the measured value in Fig. 2-3 is too negative by 7]n

according to Eq. (2-33) and correspondingly is too positive in the case of the anodic current. The error can be calculated for uniform current flow lines from Ohm's Law:

where s is the distance of the probe from the electrode surface. In the laboratory, potential measuring probes can be used and 7]Q from Eq. (2-34) can be kept very small. However, generally this is not possible for technical structures, and particu- larly not for buried objects. Possible ways to eliminate Tfa errors (i.e., by 7/?-free potential measurements) are described in Section 3.3.1.

A simplifying assumption is made that only one electrochemical reaction occurs at Ej. Then the equilibrium potential If is present at I = 0. Positive (or negative) currents can only flow with positive (or negative) deviations of if. The difference (U - if) = T] is termed the overvoltage. The function J(rj) gives infor- mation on kinetics of the reaction and on the rate-determining step. If transport through the phase boundary itself is rate determining, then /(?]) is an exponential function (activation polarization). For this reason J(U) curves are mostly plotted on a semilogarithmic scale. On the other hand, if a chemical reaction or diffusion in the medium is rate determining, then/is independent of potential, i.e., the curve 7(?7) ends parallel to the potential axis (concentration overvoltage). A similar case can arise if less readily soluble surface films are formed which in the stationary state have a solubility rate equivalent to /. This is the case with passive metals (see Section 2.3.1.2). With poorly conducting surface films or in high-resistance media, the ohmic resistance controls the current so that 7(77) follows Ohm's Law. In this case 77 is not a genuine overvoltage, but corresponds to rj^'m Eq. (2-34) and is thus in principle a measurement error, given that the potential is defined for the inter- face between the metal and film.

The literature [14] on electrochemical kinetics is extensive and specialized.

Figure 2-4 shows a/(r]) curve of a redox reaction according to Eq. (2-9) with acti- vation and diffusion polarization. It follows from theory [4, 10] for this example:

where J0 is exchange current density, and corresponds to the magnitude of the equally fast forward and reverse reactions in the equilibrium; GA and Gc are the limiting diffusion current densities and are proportional to the concentration of the reactants concerned and increase according to the first law of diffusion with the flow velocity; and 6+ and 8_ are the anodic and cathodic Tafel slopes (see definition in list of symbols) and are given by:

The term a is a symmetry factor for the energy threshold for the passage of elec- trons and is approximately equal to 0.5. In Fig. 2-4, the value of a was chosen as ²/3 for better distinction; integer exponents are chosen for 70, GA and Gc for clarity,

Fig. 2-4 Current-density-potential curves for an electrochemical partial reaction as in Eq. (2-35).

and T] is plotted dimensionless. The slopes of the curves in semilogarithmic plots (natural logarithms)

are given in the regions of Tafel lines^_ andfl+. The Tafel slopes in logs to the base 10 are obtained by multiplying by the factor In 10 = 2.303 for the system b_ and b+. At equilibrium we obtain N —> 0 and in the limiting current region, N —> <». There are similar curves for electrolytic corrosion according to Eq. (2-21) as those in Fig. 2-4. An anodic limiting current is only possible with film formation.

In general, J(U) curves for the anodic partial reaction follow a Tafel straight line. In neutral media, the cathodic partial reaction is mostly the reduction of oxy- gen, whose J(U) curve ends in a limiting cathodic current density determined by the transport of oxygen. If at high overvoltages cathodic hydrogen evolves accord- ing to Eq. (2-19), the J(U) curve bends with another constant slope.

The time dependence of the changes in the measured values is important in determining J(U) curves. In the region of the Tafel lines, stationary states are reached